It will readily be seen that if a substance which yields an ion in common with the precipitated compound is added to such a solution as has just been described, the concentration of that ion is increased, and as a result the concentration of the other ion must proportionately decrease, which can only occur through the formation of some of the undissociated compound which must separate from the already saturated solution. This explains why the addition of an excess of the precipitant is often advantageous in quantitative procedures. Such a case is discussed at length in Note 2 on page 113.

Similarly, the ionization of a specific substance in solution tends to diminish on the addition of another substance with a common ion, as, for instance, the addition of hydrochloric acid to a solution of hydrogen sulphide. Hydrogen sulphide is a weak acid, and the concentration of the hydrogen ions in its aqueous solutions is very small. The equilibrium in such a solution may be represented as:

(!(Conc'n H^{+})^{2} x Conc'n S^{—})/Conc'n H{2}S = Constant!, and a marked increase in the concentration of the H^{+} ions, such as would result from the addition of even a small amount of the highly ionized hydrochloric acid, displaces the point of equilibrium and some of the S^{—} ions unite with H^{+} ions to form undissociated H{2}S. This is of much importance in studying the reactions in which hydrogen sulphide is employed, as in qualitative analysis. By a parallel course of reasoning it will be seen that the addition of a salt of a weak acid or base to solutions of that acid or base make it, in effect, still weaker because they decrease its percentage ionization.

To understand the changes which occur when solids are dissolved where chemical action is involved, it should be remembered that no substance is completely insoluble in water, and that those products of a chemical change which are least dissociated will first form. Consider, for example, the action of hydrochloric acid upon magnesium hydroxide. The minute quantity of dissolved hydroxide dissociates thus: Mg(OH){2} Mg^{} + 2OH^{-}. When the acid is introduced, the H^{} ions of the acid unite with the OH^{-} ions to form undissociated water. The concentration of the OH^{-} ions is thus diminished, more Mg(OH){2} dissociates, the solution is no longer saturated with the undissociated compound, and more of the solid dissolves. This process repeats itself with great rapidity until, if sufficient acid is present, the solid passes completely into solution.

Exactly the same sort of process takes place if calcium oxalate, for example, is dissolved in hydrochloric acid. The C_{2}O_{4}^{—} ions unite with the H^{} ions to form undissociated oxalic acid, the acid being less dissociated than normally in the presence of the H^{} ions from the hydrochloric acid (see statements regarding hydrogen sulphide above). As the undissociated oxalic acid forms, the concentration of the C_{2}O_{4}^{—} ions lessens and more CaC_{2}O_{4} dissolves, as described for the Mg(OH)_{2} above. Numerous instances of the applications of these principles are given in the Notes.

Water itself is slightly dissociated, and although the resulting H^{+} and OH^{-} ions are present only in minute concentrations (1 mol. of dissociated water in 10^{7} liters), yet under some conditions they may give rise to important consequences. The term !hydrolysis! is applied to the changes which result from the reaction of these ions. Any salt which is derived from a weak base or a weak acid (or both) is subject to hydrolytic action. Potassium cyanide, for example, when dissolved in water gives an alkaline solution because some of the H^{+} ions from the water unite with CN^{-} ions to form (HCN), which is a very weak acid, and is but very slightly dissociated. Potassium hydroxide, which might form from the OH^{-} ions, is so largely dissociated that the OH^{-} ions remain as such in the solution. The union of the H^{+} ions with the CN^{-} ions to form the undissociated HCN diminishes the concentration of the H^{+} ions, and more water dissociates (H{2}O H^{+} + OH^{-}) to restore the equilibrium. It is clear, however, that there must be a gradual accumulation of OH^{-} ions in the solution as a result of these changes, causing the solution to exhibit an alkaline reaction, and also that ultimately the further dissociation of the water will be checked by the presence of these ions, just as the dissociation of the H{2}S was lessened by the addition of HCl.

An exactly opposite result follows the solution of such a salt as Al{2}(SO{4}){3} in water. In this case the acid is strong and the base weak, and the OH^{-} ions form the little dissociated Al(OH){3}, while the H^{+} ions remain as such in the solution, sulphuric acid being extensively dissociated. The solution exhibits an acid reaction.

Such hydrolytic processes as the above are of great importance in analytical chemistry, especially in the understanding of the action of indicators in volumetric analysis. (See page 32.)

The impelling force which causes an element to pass from the atomic to the ionic condition is termed !electrolytic solution pressure!, or ionization tension. This force may be measured in terms of electrical potential, and the table below shows the relative values for a number of elements.

In general, an element with a greater solution pressure tends to cause the deposition of an element of less solution pressure when placed in a solution of its salt, as, for instance, when a strip of zinc or iron is placed in a solution of a copper salt, with the resulting precipitation of metallic copper.

Hydrogen is included in the table, and its position should be noted with reference to the other common elements. For a more extended discussion of this topic the student should refer to other treatises.

POTENTIAL SERIES OF THE METALS

POTENTIAL POTENTIAL IN VOLTS IN VOLTS Sodium Na^{+} +2.44 Lead Pb^{} -0.13 Calcium Ca^{} Hydrogen H^{} -0.28 Magnesium Mg^{} Bismuth Bi^{} Aluminum A1^{} 1.00 Antimony -0.75 Manganese Mn^{} Arsenic Zinc Zn^{+} +0.49 Copper Cu^{} -0.61 Cadmium Cd^{} 0.14 Mercury Hg^{} -1.03 Iron Fe^{+} +0.063 Silver Ag^{} -1.05 Cobalt Co^{} -0.045 Platinum Nickel Ni^{} -0.049 Gold Tin Sn^{} -0.085(?)



THE FOLDING OF A FILTER PAPER

If a filter paper is folded along its diameter, and again folded along the radius at right angles to the original fold, a cone is formed on opening, the angle of which is 60 deg.. Funnels for analytical use are supposed to have the same angle, but are rarely accurate. It is possible, however, with care, to fit a filter thus folded into a funnel in such a way as to prevent air from passing down between the paper and the funnel to break the column of liquid in the stem, which aids greatly, by its gentle suction, in promoting the rate of filtration.

Such a filter has, however, the disadvantage that there are three thicknesses of paper back of half of its filtering surface, as a consequence of which one half of a precipitate washes or drains more slowly. Much time may be saved in the aggregate by learning to fold a filter in such a way as to improve its effective filtering surface. The directions which follow, though apparently complicated on first reading, are easily applied and easily remembered. Use a 6-inch filter for practice. Place four dots on the filter, two each on diameters which are at right angles to each other. Then proceed as follows: (1) Fold the filter evenly across one of the diameters, creasing it carefully; (2) open the paper, turn it over, rotate it 90 deg. to the right, bring the edges together and crease along the other diameter; (3) open, and rotate 45 deg. to the right, bring edges together, and crease evenly; (4) open, and rotate 90 deg. to the right, and crease evenly; (5) open, turn the filter over, rotate 22-(1/2) deg. to the right, and crease evenly; (6) open, rotate 45 deg. to the right and crease evenly; (7) open, rotate 45 deg. to the right and crease evenly; (8) open, rotate 45 deg. to the right and crease evenly; (9) open the filter, and, starting with one of the dots between thumb and forefinger of the right hand, fold the second crease to the left over on it, and do the same with each of the other dots. Place it, thus folded, in the funnel, moisten it, and fit to the side of the funnel. The filter will then have four short segments where there are three thicknesses and four where there is one thickness, but the latter are evenly distributed around its circumference, thus greatly aiding the passage of liquids through the paper and hastening both filtration and washing of the whole contents of the filter.

!SAMPLE PAGES FOR LABORATORY RECORDS!

!Page A!

Date

CALIBRATION OF BURETTE No.

_____________ BURETTE DIFFERENCE OBSERVED DIFFERENCE CALCULATED READINGS WEIGHTS CORRECTION ___ ___ ___ ___ ___ 0.02 16.27 10.12 10.10 26.35 10.08 -.02 20.09 9.97 36.26 9.91 -.06 30.16 10.07 46.34 10.08 +.01 40.19 10.03 56.31 9.97 -.06 50.00 9.81 66.17 9.86 +.05 ___ ___ ___ ___ ___

These data to be obtained in duplicate for each burette.

!Page B!

Date

DETERMINATION OF COMPARATIVE STRENGTH HCl vs. NaOH

_____________ DETERMINATION I II _____ ____ ____ Corrected Corrected Final Reading HCl 48.17 48.08 43.20 43.14 Initial Reading HCl 0.12 .12 .17 .17 - - - - 47.96 42.97 Corrected Corrected Final Reading HCl 46.36 46.29 40.51 40.37 Initial Reading HCl 1.75 1.75 .50 .50 - - - - 44.54 39.87 log cc. NaOH 1.6468 1.6008 colog cc. HCl 8.3192 8.3668 9.9680 - 10 9.9676 - 10 1 cc. HCl .9290 cc. NaOH .9282 cc. NaOH Mean .9286 _____ ____ ____

Signed

!Page C! Date

STANDARDIZATION OF HYDROCHLORIC ACID ===================================================================== Weight sample and tube 9.1793 8.1731 8.1731 6.9187 Weight sample 1.0062 1.2544 Final Reading HCl 39.97 39.83 49.90 49.77 Initial Reading HCl .00 .00 .04 .04 - - - - 39.83 49.73 Final Reading NaOH .26 .26 .67 .67 Initial Reading NaOH .12 .12 .36 .36 - - - - .14 .31 .14 .31 Corrected cc. HCl 39.83 - - = 39.68 49.73 - - = 49.40 .9286 .9286 log sample 0.0025 0.0983 colog cc 8.4014 - 10 8.3063 - 10 colog milli equivalent 1.2757 1.2757 9.6796 - 10 9.6803 - 10 Normal value HCl .4782 .4789 Mean .4786 =====================================================================

Signed

!Page D! Date

DETERMINATION OF CHLORINE IN CHLORIDE, SAMPLE No. ===================================================================== Weight sample and tube 16.1721 15.9976 15.9976 15.7117 - - Weight sample .1745 .2859 Weight crucible + precipitate 14.4496 15.6915 Constant weights 14.4487 15.6915 14.4485 Weight crucible 14.2216 15.3196 Constant weight 14.2216 15.3194 Weight AgCl .2269 .3721 log Cl 1.5496 1.5496 log weight AgCl 9.3558 - 10 9.5706 - 10 log 100 2.0000 2.0000 colog AgCl 7.8438 - 10 7.7438 - 10 colog sample 0.7583 0.5438 - - 1.5075 1.5078 Cl in sample No. 32.18% 32.20% =====================================================================

Signed

STRENGTH OF REAGENTS

The concentrations given in this table are those suggested for use in the procedures described in the foregoing pages. It is obvious, however, that an exact adherence to these quantities is not essential.

Approx. Approx. Grams relation relation per to normal to molal liter. solution solution

Ammonium oxalate, (NH_{4})_{2}C_{2}O_{4}.H_{2}O 40 0.5N 0.25 Barium chloride, BaCl_{2}.2H_{2}O 25 0.2N 0.1 Magnesium ammonium chloride (of MgCl_{2}) 71 1.5N 0.75 Mercuric chloride, HgCl_{2} 45 0.33N 0.66 Potassium hydroxide, KOH (sp. gr. 1.27) 480 Potassium thiocyanate, KSCN 5 0.05N 0.55 Silver nitrate, AgNO_{3} 21 0.125N 0.125 Sodium hydroxide, NaOH 100 2.5N 2.5 Sodium carbonate. Na_{2}CO_{3} 159 3N 1.5 Sodium phosphate, Na_{2}HPO_{4}.12H_{2}O 90 0.5N or 0.75N 0.25

Stannous chloride, SnCl_{2}, made by saturating hydrochloric acid (sp. gr. 1.2) with tin, diluting with an equal volume of water, and adding a slight excess of acid from time to time. A strip of metallic tin is kept in the bottle.

A solution of ammonium molybdate is best prepared as follows: Stir 100 grams of molybdic acid (MoO_{3}) into 400 cc. of cold, distilled water. Add 80 cc. of concentrated ammonium hydroxide (sp. gr. 0.90). Filter, and pour the filtrate slowly, with constant stirring, into a mixture of 400 cc. concentrated nitric acid (sp. gr. 1.42) and 600 cc. of water. Add to the mixture about 0.05 gram of microcosmic salt. Filter, after allowing the whole to stand for 24 hours.

The following data regarding the common acids and aqueous ammonia are based upon percentages given in the Standard Tables of the Manufacturing Chemists' Association of the United States [!J.S.C.I.!, 24 (1905), 787-790]. All gravities are taken at 15.5 deg.C. and compared with water at the same temperature.

Aqueous ammonia (sp. gr. 0.96) contains 9.91 per cent NH_{3} by weight, and corresponds to a 5.6 N and 5.6 molal solution.

Aqueous ammonia (sp. gr. 0.90) contains 28.52 per cent NH_{3} by weight, and corresponds to a 5.6 N and 5.6 molal solution.

Hydrochloric acid (sp. gr. 1.12) contains 23.81 per cent HCl by weight, and corresponds to a 7.3 N and 7.3 molal solution.

Hydrochloric acid (sp. gr. 1.20) contains 39.80 per cent HCl by weight, and corresponds to a 13.1 N and 13.1 molal solution.

Nitric acid (sp. gr. 1.20) contains 32.25 per cent HNO_{3} by weight, and corresponds to a 6.1 N and 6.1 molal solution:

Nitric acid (sp. gr. 1.42) contains 69.96 per cent HNO_{3} by weight, and corresponds to a 15.8 N and 15.8 molal solution.

Sulphuric acid (sp. gr. 1.8354) contains 93.19 per cent H{2}SO{4} by weight, and corresponds to a 34.8 N or 17.4 molal solution.

Sulphuric acid (sp. gr. 1.18) contains 24.74 per cent H{2}SO{4} by weight, and corresponds to a 5.9 N or 2.95 molal solution.

The term !normal! (N), as used above, has the same significance as in volumetric analyses. The molal solution is assumed to contain one molecular weight in grams in a liter of solution.

DENSITIES AND VOLUMES OF WATER AT TEMPERATURES FROM 15-30 deg.C.

Temperature Density. Volume. Centigrade.

4 deg. 1.000000 1.000000 15 deg. 0.999126 1.000874 16 deg. 0.998970 1.001031 17 deg. 0.998801 1.001200 18 deg. 0.998622 1.001380 19 deg. 0.998432 1.001571 20 deg. 0.998230 1.001773 21 deg. 0.998019 1.001985 22 deg. 0.997797 1.002208 23 deg. 0.997565 1.002441 24 deg. 0.997323 1.002685 25 deg. 0.997071 1.002938 26 deg. 0.996810 1.003201 27 deg. 0.996539 1.003473 28 deg. 0.996259 1.003755 29 deg. 0.995971 1.004046 30 deg. 0.995673 1.004346

Authority: Landolt, Boernstein, and Meyerhoffer's !Tabellen!, third edition.

CORRECTIONS FOR CHANGE OF TEMPERATURE OF STANDARD SOLUTIONS

The values below are average values computed from data relating to a considerable number of solutions. They are sufficiently accurate for use in chemical analyses, except in the comparatively few cases where the highest attainable accuracy is demanded in chemical investigations. The expansion coefficients should then be carefully determined for the solutions employed. For a compilation of the existing data, consult Landolt, Boernstein, and Meyerhoffer's !Tabellen!, third edition.

Corrections for 1 cc. Concentration. of solution between 15 deg. and 35 deg.C.

Normal .00029 0.5 Normal .00025 0.1 Normal or more dilute solutions .00020

The volume of solution used should be multiplied by the values given, and that product multiplied by the number of degrees which the temperature of the solution varies from the standard temperature selected for the laboratory. The total correction thus found is subtracted from the observed burette reading if the temperature is higher than the standard, or added, if it is lower. Corrections are not usually necessary for variations of temperature of 2 deg.C. or less.



INTERNATIONAL ATOMIC WEIGHTS

========================================================== 1920 1920 Aluminium Al 27.1 Molybdenum Mo 96.0 Antimony Sb 120.2 Neodymium Nd 144.3 Argon A 39.9 Neon Ne 20.2 Arsenic As 74.96 Nickel Ni 58.68 Barium Ba 137.37 Nitrogen N 14.008 Bismuth Bi 208.0 Osmium Os 190.9 Boron B 11.0 Oxygen O 16.00 Bromine Br 79.92 Palladium Pd 106.7 Cadmium Cd 112.40 Phosphorus P 31.04 Caesium Cs 132.81 Platinum Pt 195.2 Calcium Ca 40.07 Potassium K 39.10 Carbon C 12.005 Praseodymium Pr 140.9 Cerium Ce 140.25 Radium Ra 226.0 Chlorine Cl 35.46 Rhodium Rh 102.9 Chromium Cr 52.0 Rubidium Rb 85.45 Cobalt Co 58.97 Ruthenium Ru 101.7 Columbium Cb 93.1 Samarium Sm 150.4 Copper Cu 63.57 Scandium Sc 44.1 Dysprosium Dy 162.5 Selenium Se 79.2 Erbium Er 167.7 Silicon Si 28.3 Europium Eu 152.0 Silver Ag 107.88 Fluorine Fl 19.0 Sodium Na 23.00 Gadolinium Gd 157.3 Strontium Sr 87.63 Gallium Ga 69.9 Sulphur S 32.06 Germanium Ge 72.5 Tantalum Ta 181.5 Glucinum Gl 9.1 Tellurium Te 127.5 Gold Au 197.2 Terbium Tb 159.2 Helium He 4.00 Thallium Tl 204.0 Hydrogen H 1.008 Thorium Th 232.4 Indium In 114.8 Thulium Tm 168.5 Iodine I 126.92 Tin Sn 118.7 Iridium Ir 193.1 Titanium Ti 48.1 Iron Fe 55.84 Tungsten W 184.0 Krypton Kr 82.92 Uranium U 238.2 Lanthanum La 139.0 Vanadium V 51.0 Lead Pb 207.2 Xenon Xe 130.2 Lithium Li 6.94 Ytterbium Yb 173.5 Lutecium Lu 175.0 Yttrium Y 88.7 Magnesium Mg 24.32 Zinc Zn 65.37 Manganese Mn 54.93 Zirconium Zr 90.6 Mercury Hg 200.6 ==========================================================



INDEX

Acidimetry Acid solutions, normal standard Acids, definition of Acids, weak, action of other acids on action of salts on Accuracy demanded Alkalimetry Alkali solutions, normal standard Alumina, determination of in stibnite Ammonium nitrate, acid Analytical chemistry, subdivisions of Antimony, determination of, in stibnite Apatite, analysis of Asbestos filters Atomic weights, table of

Balances, essential features of use and care of Barium sulphate, determination of sulphur in Bases, definition of Bichromate process for iron Bleaching powder, analysis of Brass, analysis of Burette, description of calibration of cleaning of reading of

Calcium, determination of, in limestone Calibration, definition of of burettes of flasks Carbon dioxide, determination of, in limestone Chlorimetry Chlorine, gravimetric determination of Chrome iron ore, analysis of Coin, determination of silver in Colloidal solution of precipitates Colorimetric analyses, definition of Copper, determination of, in brass determination of in copper ores Crucibles, use of Crystalline precipitates

Densities of water Deposition potentials Desiccators Direct methods Dissociation, degree of

Economy of time Electrolytic dissociation, theory of Electrolytic separations, principles of End-point, definition of Equilibrium, chemical Evaporation of liquids

Faraday's law Feldspar, analysis of Ferrous ammonium sulphate, analysis of Filters, folding of how fitted Filtrates, testing of Filtration Flasks, graduation of Funnels Fusions, removal of from crucibles

General directions for gravimetric analysis volumetric analysis Gooch filter Gravimetric analysis, definition of

Hydrochloric acid, standardization of Hydrolysis

Ignition of precipitates Indicators, definition of for acidimetry preparation of Indirect methods Insoluble matter, determination of in limestone Integrity Iodimetry Ions, definition of Iron, gravimetric determination of volumetric determination of

Jones reductor

Lead, determination of in brass Limestone, analysis of Limonite, determination of iron in Liquids, evaporation of transfer of Litmus Logarithms

Magnesium, determination of Mass action, law of Measuring instruments Methyl orange Moisture, determination of in limestone

Neutralization methods Normal solutions, acid and alkali oxidizing agents reducing agents Notebooks, sample pages of

Oxalic acid, determination of strength of Oxidation processes Oxidizing power of pyrolusite

Permanganate process for iron Phenolphthalein Phosphoric anhydride, determination of Pipette, calibration of description of Platinum crucibles, care of Precipitates, colloidal crystalline ignition of separation from filter washing of Precipitation Precipitation methods (volumetric) Problems Pyrolusite, oxidizing power of

Quantitative Analyses, subdivisions of

Reagents, strength of Reducing solution, normal Reductor, Jones Reversible reactions

Silica, determination of, in limestone determination of, in silicates purification of Silicic acid, dehydration of Silver, determination of in coin Soda ash, alkaline strength of Sodium chloride, determination of chlorine in Solubility product Solution pressure Solutions, normal standard Standardization, definition of Standard solutions, acidimetry and alkalimetry chlorimetry iodimetry oxidizing and reducing agents thiocyanate Starch solutions Stibnite, determination of antimony in Stirring rods Stoichiometry Strength of reagents Suction, use of Sulphur, determination of in ferrous ammonium sulphate in barium sulphate

Temperature, corrections for Testing of washings Theory of electrolytic dissociation Thiocyanate process for silver Titration, definition of Transfer of liquids

Volumetric analysis, definition of general directions

Wash-bottles Washed filters Washing of precipitates Washings, testing of Water, ionization of densities of Weights, care of

Zimmermann-Reinhardt method for iron Zinc, determination of, in brass

THE END